Molecular and Ionic Compound Structure and Properties

1. Molecular and Ionic Compound Structure and Properties AP Chemistry

2. Topic 2.1 Types of Chemical Bonds BL 8.1 Problems: 8.6

3. Topic 2.1 Types of Chemical Bonds BL 8.1, # 8.6 Chemical bonds hold atoms together. There are 3 types of chemical bonds: Ionicbonds(electrostatic forces that hold ions together…) Example: Na+Cl-, K+Br- Covalentbonds(result from sharing electrons between atoms…) Example: H2, NH3 Metallicbonds (refers to metal nuclei floating in a sea of electrons…) Example: copper, gold

4. Lewis Symbols: Electron dots Valence Electrons: outer-most electrons determined by the “Group A #” from the periodic table Exceptions: d or f-block = 2 valence electrons & Helium =2 e-

5. Octet Rule Atoms often gain or lose or share electrons to fill their valence shell with 8 electrons to achieve a noble gas configuration. Exceptions: * Hydrogen needs only 2 e- to be filled. * Some nonmetals can have more or less than 8.

6. Ionic Bonding Topic 2.1 and 2.3 BL 8.2, 11.8, #’s 12,16,20,70 Problems: 12, 16, 20, 70

7. Ionic Bonding Topic 2.1 and 2.3 BL 8.2, 11.8 • Ionicbonds—transfer of electrons • Form between an element of low ionizationenergy (not much energy required to pull off an electron) and an element of high electron affinity (lots of energy is released when an electron is added to its outer shell). • Usually form between a metal and a nonmetal. • The best way to determine if a compound is ionic or covalent is to examine the properties

8. Ionic Bonding Energies Consider the reaction between sodium and chlorine: Na(s) + ½Cl2(g) NaCl(s)∆Hºf = –410.9 kJ/mol - The reaction is violently exothermic. - We infer that the NaCl is more stable than its constituent elements. Here’s another way to look at the energy of ionic bond formation: Sodium loses 1 electron…Na  Na+ + 1 e- Requires 5.1 eV of energy Chlorine gains l electron…Cl + 1 e-  Cl- Releases 3.6 eV of energy NaCl forms… Na+ + Cl- [Na+][ Cl-] Releases 5.2 eV of energy [1 eV (electron volt) = 1.602 x 10-19 J] The energy released is greater than the energy required, therefore the ionic bond forms… (∆Hf = - 3.7 eV)

9. Lattice Energy Lattice energy is the energy required to completely separate a mole of a solid ionic compound into its gaseous ions. Lattice energy increases as distance between the ions decreases. Lattice energy increases as charges on the ions increase.

10. Ions Metals lose electrons to form smaller (+) cations. Nonmetals gain e- to form larger (-) anions. The # of e- gained or lost depends on how many they need to gain or lose to get to a noble gas configuration. Only then will they become stable. Groups of atoms can have charges too. They are called polyatomic ions. The atoms share electrons (covalent bonds) but the group still has an overall charge. Examples: [NH4+] , [CO3-2]

11. Metallic Compounds Topic 2.1 and 2.4 BL 11.8, 23.5, 23.6, #’s 23.20, 27, 28 Problems: 23.20, 27, 28

12. Metallic Compounds Topic 2.1 and 2.4 BL 11.8, 23.5, 23.6 • Metallic bond: Electrostatic force of attraction between metal ions and mobile valence electrons • Referred to as the electron-sea model • Due to these mobile electrons metals are • Malleable • Ducticle • High conductivity of heat and electricity

13. Metallic Bonding

14. Types of Metal Alloys • Substitutional: formed when 2 metallic components have similar atomic radii and chemical bonding characteristics • Example: brass alloys • Intersitial: smaller atoms will occupy holes between metal atoms • Example: Steel

15. Types of Alloys

16. Topic 2.1 and 2.2 Covalent Bonds BL 8.3,8.4 and 9.4, #’s 8.26, 30, 34, 36 Problems: 8.26, 30, 34, and 36

17. Topic 2.1 and 2.2 Covalent Bonds BL 8.3,8.4 and 9.4 - A line can also be used to represent 2 shared e-’s (or one covalent bond.) Atoms share electrons to fill their valence shell. Usually form between 2 nonmetals Lewis Structures: represent covalent bonds as 2 dots between the atoms

18. Bonds form when orbitals on atoms overlap. • There are two electrons of opposite spin in the overlapping orbitals. Why do bonds form?

19. The overlapping of the orbitals will lower the overall energy of the 2 atoms, therefore it is more stable. Why do bonds form?

20. Multiple Covalent Bonds Single bond = 2 electrons shared …(1 pair) Double bond= 4 electrons shared…(2 pairs) Triple bond= 6 electrons shared…(3 pairs) Bond Lengths & Bond Strengths Single bonds are the longest and weakest covalent bonds. Triple are the shortest and strongest covalent bonds. Ionic bonds are much stronger than covalent bonds.

21. Bond Polarity Bond polarity helps to describe the sharing of the electrons between atoms. There are 3 possibilities… Nonpolar covalent: equal sharing of the e- pair Polar covalent: unequal sharing of the e- pair Ionic: transfer of valence e- from the metal to the nonmetal A molecule that has one side slightly positive and one side slightly negative is said to be a “dipole.” The positive end (or pole) in a polar bond is represented + and the negative pole -. Arrow can also show dipoles. + -

22. Bond Polarity & Electronegativity (How can you tell what type of bond will form?) Electronegativity: describes an atom’s attraction to the e- pair in a bond…(It’s a number from 0 to 4.0) The difference between electronegativities indicate whether a bond will be nonpolar, polar or ionic. There is no sharp distinction between bonding types. In General: Nonpolar = 0-0.4 Polar= 0.5-2.0 Ionic= Above 2.0 Lattice Energy and Polarity Lattice energy increases as the electronegativity between the atoms in an ionic compound increases

23. Dipole Moments • Consider HF: • The difference in electronegativity leads to a polar bond. • There is more electron density on F than on H. • Since there are two different “ends” of the molecule, we call HF a dipole. • Dipole moment(debyes, D), m, is the magnitude of the dipole: • µ=Qr • where Q is the magnitude of the charges. Chapter 8

24. Topic 2.5Rules For Drawing Lewis Structures BL 8.5 Problem: 8.44

25. Rule 1) How many electrons are possible around an atom? • For hydrogen, only 2 electrons are possible, therefore only one bond! • Second row elements usually try to get 8 e-. Notable Exception: Boron needs only 6. • Third & Fourth Row usually have 8 e- but can expand to get 10 or more. Rule 2) Drawing the Lewis Structure: • First arrange the atoms around the central atom (usually the highest electronegative one, but never hydrogen!) • Count the total # of valence electrons in the molecule. If it is an ion, add 1 for each (-) charge or subtract 1 for each (+) charge. • Distribute the electrons keeping Rule #1 in mind. If you have too many electrons, look for double or triple bonds, or place the extras around the “3rd or 4th row element.” Topic 2.5Rules For Drawing Lewis Structures BL 8.5

26. Rule 3) For Ionic Compounds: • Indicate the charge on the ions • Do not share the electrons, transfer them! Rule 4) For odd-numbered valence electrons: • If you must cheat an element out of 8 e- and only give it 7, then cheat the least electronegative element. Rule 5) Resonance Structures: • If there is more than one way to draw the Lewis structure, show them all. Rule 6) Nature doesn’t know anything about Rules1-5. • These are just models for us to use in order to represent bonding. Rules For Drawing Lewis Structures

27. Drawing Lewis Structures • Sum valence electron from all atoms. • Write symbols for the atoms and show which atoms are connected to which. • Complete the octets of the atoms bonded to the central atom. • Place leftover electrons on the central atom, even if it violates the octet rule. • If there are not enough electrons to give the central atom an octet, try multiple bonds. Chapter 8

28. Topic 2.6Resonance BL 8.6, 8.7, # 8.46, 52, 80 Problems: 8.46, 52, 80

29. Resonance Structures: • The ability to draw more than one “correct” Lewis structure. Topic 2.6Resonance BL 8.6, 8.7 Benzene (C6H6) • The true structure for the molecule is somewhere “in between” the resonance structures.

30. Formal Charge: The formal charge of an atom is the charge that an atom (in a molecule) would have if all of the atoms had the same electronegativity. • To calculate formal charge: (valence e- - # of bonds - lone pair e-) Practice: Determine the formal charge on C and N. Formal Charge C= 4 – 3 – 2 = -1 N= 5 – 3 – 2 = 0

31. The most stable Lewis structure has the smallest formal charge on each atom and the most negative formal charge on the most electronegative atoms. • It is important to keep in mind that formal charges do NOT represent REAL charges on atoms! Formal Charge

32. Topic 2.7 Molecular Geometry & Bonding Theories, BL 9.1-3, 9.5, 9.6 Problems: 9.6,10,12,16,18,20, 26, 28, 32, 38, 44, 46, 58, 73, 76, 83

33. Topic 2.7 Molecular Geometry & Bonding Theories, BL 9.1-3, 9.5, 9.6 Lewis structures tell us which atoms are bonded together, but we will now explore the geometric shapes of these molecules. Overall shape is determined by bond angles. Bond angles are determined by the VSEPR theory. Electrons repel & will try to get as far away from each other as possible Nonbonded electron pairs take up more space than bonded electrons. We will first explore the molecular form ABn. This represents one central atom bonded to 1-6 other atoms. First you must determine the # of electron domains on the central atom. An electron domain is a region of electrons that are either bonded or non-bonded (lone pairs). A double or triple bond only counts as one domain.

34. Electron Domain Geometry The arrangement of electron domains about the central atom of an ABn molecule is its electron-domain geometry. There are five different electron-domain geometries: linear --(2 electron domains) trigonal planar --(3 domains) tetrahedral --(4 domains) trigonal bipyramidal --(5 domains) octahedral --(6 domains).

35. Chapter 9

36. Electron Domain Geometry

37. For example… :O=C=O: • There are 2 electron domains on carbon…Its shape must therefore be linear. H–O–H • There are 4 electron domains on oxygen….Its shape is based on the tetrahedral. Electron Domain Geometry .. ..

38. The molecular geometry is the arrangement of the atoms in space. To determine the shape of a molecule we will distinguish between bonding pairs and lone pairs. • Count the # of bonding domains vs. nonbonding domains. H-O-H Oxygen has 2 bonding and 2 nonbonding domains • With this information, we can determine the molecular geometry…bent (as we know already!) Molecular Geometry .. ..

39. Molecular Geometry

40. Molecular Geometry

41. Molecular Geometry

42. The most common shapes we deal with are as follows: • Tetrahedral, pyramidal, bent, linear, and trigonal planar. (It is to your advantage to know some common examples of each of these shapes!!) • The “ideal” bond angle between the central atom and the other atoms should be noted… Linear= 180º Tetrahedral = 109.5º Trigonal Planar =120º • Due to the lone pairs of electrons on pyramidal and bent shapes, the ideal bond angles will be less than 109.5º Molecular Geometry—Most Common Shapes

43. In general, multiple bonds repel more as do lone pairs. Molecular Geometry— e- repulsion

44. In acetic acid, CH3COOH, there are three interior atoms: two C’s and one O. •We assign the molecular (and electron-domain) geometry about each interior atom separately: -The geometry around the first C is tetrahedral. -The geometry around the second C is trigonal planar. -The geometry around the O is bent (tetrahedral). Shapes of Larger Molecules

45. When there is a difference in electronegativity between two atoms, then the bond between them is polar. • It is possible for a molecule to contain polar bonds, but not be polar. -For example, the bond dipoles in CO2 cancel each other because CO2 is linear. Molecular Shape and Molecular Polarity

46. In water, the molecule is not linear and the bond dipoles do not cancel each other. Therefore, water is a polar molecule. Molecular Shape and Molecular Polarity

47. The overall polarity of a molecule depends on its molecular geometry. Molecular Shape and Molecular Polarity

48. A hybrid orbital is simply a mixing of different orbitals together to form a new “hybridized orbital”. • We need the concept of hybrid orbitals to explain molecular shapes. (Let’s try to keep it simple…) When you mix n atomic orbitals we must get n hybrid orbitals. Hybrid Orbitals Example: If you mix one “s” orbital and three “p” orbitals you will get four “sp3” hybrid orbitals that all have exactly the same energies.

49. Hybrid Orbitals The # of electron domains on the atom will indicate the hybridization needed. Example: H2C=CH2 (Carbon has 3 e- domains so its hybridization must be sp2 which has 3 hybrid orbital domains as well.)

50. Overlapping orbitals come in 2 varieties… • -Bonds: electron density lies on the axis between the nuclei. - All single bonds are -bonds. Sigma and Pi Bonds